Equilibrium

1. Equilibrium to what extent will the reaction proceed

2. Solutions • Solubility -the amount of solute that can be dissolved to form a solution. • Solvent – the substance in a solution present in the greatest amount. • Solute – the substance in a solution present in the least amount. • Saturate – a solution that has come to equilibrium. The rate of dissolving is equal to the rate of recrystalizing. • High solubility – more than .10 mol/L will dissolve (AgNO3) • Low solubility – less than .10 mol/L will dissolve (AgCl)

3. Dissociation (ionization) Reactions When ionic compounds dissolve in water, they separate into ions – one positive and one negative. Examples: Ca(NO3)2(s)→ Ca2+(aq) + 2 NO3–(aq) Ca3(PO4)2→ 3 Ca2+(aq) + 2 PO43-(aq) Write dissociation reactions for each of the following: K2SO4(s)→ Al(OH)3(s)→ PbSO4(s)→ NH4NO3(s) → 2 K+(aq) + SO4 2-(aq) Al3+(aq) + 3 OH –(aq) (slightly soluble) Pb2+(aq) + SO42-(aq) (slightly soluble) NH4+(aq) + NO3–(aq)

4. net ionic equation Net Ionic Equations Combine solutions of NaCl(aq) and AgNO3(aq) NaCl(aq) + AgNO3(aq)→ AgCl(s) + NaNO3(aq) Na+(aq) + Cl-(aq) + Ag+(aq) + NO3–(aq)→ AgCl(s) +Na+(aq) + NO3–(aq) Cl-(aq) + Ag+(aq)→ AgCl(s)

5. Collision Theory A mathematical description of the number of collisions between molecules in a sample of matter per unit time, useful for predicting rates of reaction.

6. Collision Theory Reaction rate: the number of atoms, ions or molecules that react in a given period of time to form products

7. THE STATES OF THE COLLISION THEORY Initial State Final State I2 H2 HI

8. HI HI Collision Theory Collision Theory states that in order for reactions to occur between substances, their particles (molecules, atoms, or ions) must collide. These interactions, if effective will form two new molecules.

9. HI HI HI HI A COLLISION THAT IS TOO GENTLE • This collision is not energetic enough to supply the required activation energy. • Therefore the Collision is ineffective. HI HI

10. HI HI HI HI A COLLISION IN POOR ORIENTATION • The colliding molecules are not oriented in a way that enables them to react with each other. • If the Collision doesn’t have the right orientation then the collision is noteffective. HI HI

11. AN EFFECTIVE COLLISION • This collision has the right orientation. • This collision is powerful enough to cause a good effect. • Everything is satisfied, and the collision turns out to be effective. HI HI HI HI I I H2

12. Activation Energy • Eact - minimum energy a reactant must possess in order to convert to products. • The activation energy (Eact) can determine how fast a reaction occurs. In general, the higher the activation energy, the slower the reaction rate. The lower the activation energy, the faster the reaction.

13. Activation Energy Consider the process of someone trying to roll a boulder over a hill. The higher the hill, the slower the task. The lower the hill the faster the process. The height of the hill (a) correspond to the energy of activation (Eact). E act

14. Activated complex • the atomic configuration at the top of the energy barrier • short life time (10-13 s) • breaks apart to form reactants or products both of which have lower potential energy activated complex

15. activated complex

16. Activation Energy • This explains why some reactions do not take place at room temperature. • CH4(g) + O2(g)→ no reaction at room temp. • The molecules can not overcome the activation energy. A reaction with a very low activation energy will occur spontaneously.

17. Types of Reactions Spontaneous exothermic Spontaneous endothermic Slow exothermic Slow endothermic

18. Reaction Mechanism • Most reactions do not take place in a single step. A→B→C→D • Each step is usually a simple one on one collision reaction. • The set of steps is called the reaction mechanism. • The slowest step in the reaction mechanism is the rate determining step.

19. Reaction Mechanism rate determining step, highest activation energy and therefore the slowest rate. Activated Complex II III Energy I Reactants Products Intermediates Reaction Progress

20. Catalyst Effect A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the reaction. 2H2(g) + O2(g) → 2H2O(l) Slow 2H2(g) + O2(g) → 2H2O(l) Fast The Pt is not used up and does not appear as a reactant or product Pt

21. Catalyst Effect Consider the task of moving coal over a barrier. A pathway with a lower barrier is analogous to a reaction affected by a catalyst. The task becomes easier for a pathway with a lower barrier .

23. Catalyst Effect • Enzymes are organic catalysts the allow chemical reactions in the human body to occur at a lower temperature than normal.

24. Catalyst Effect Inhibitors effect catalysts by rendering them useless. (a) Cholinesterase catalyzes the hydrolysis of acetylcholine into acetic acid and choline. (b) An organic phosphate (as in the nerve gas Sarin) binds to cholinesterase, preventing it from breaking down acetylcholine.

26. Catalytic Converter A catalytic converter works by taking exhaust gases from the engine, including CO and NO, passing them through the catalytic converter, where they are converted to CO2 and N2 by catalyzed reactions.

27. Reaction Conditions Affecting Rates Surface area- the more surface area, the greater the chance for reactants to encounter to form product. • Catalyst- lowers the activation energy for the reaction.

28. Temperature -the higher the temperature the faster the molecules will move: • higher frequency of collisions • more energy in each collision Nature of the reactants • solids and liquids cannot undergo a change in concentration since they occupy a given space determined by intermolecular bonding. • gases and solutions can alter their concentrations.

29. Concentration-The higher the concentration the more particles per unit area – higher probability of a collision

30. Reversible Reactions • The conversion of reactants into products and the conversion of products into reactants occur simultaneously. • 2SO2(g) + O2(g) 2SO3(g) • In a reversible reaction the rate of the reverse reaction is zero at the start.

31. Terms and Symbols • Eafwd → activation energy for the forward reaction • Earev → activation energy for the reverse reaction

33. Exothermic Endothermic Comparing Types of Reactions Earev Eafwd Eafwd Ep Ep Earev ΔH ΔH

34. Catalyst Affecting Reaction Rate • The catalyst lowers the E act for the forward and the reverse reaction because the reaction takes place through a different set of steps.

35. Equilibrium Many reactions do not convert 100% of reactants to products. There is often a point in a reaction when the products will back react to form reactants. • The extent of the reaction, 20% or 80%, can be determined by measuring the concentration of each component in solution. • In general the extent of the reaction is a function of temperature and concentration which is monitored by some constant value called the equilibrium constant (Keq).

36. Dynamic Equilibrium • Chemical Equilibrium is a dynamic state in which the rates of the forward and the reverse reaction are equal.

38. p q     P • Q Keq  a b     A • B Dynamic Equilibrium • For any general chemical process at equilibrium. aA + bB  pP + qQ eq

39. Concentration vs time graph for the reversible reaction 2Hl(g)↔ H2(g) + I2(g) After the time, te, the reaction is at equilibrium, and the concentrations of reactant and products undergo no further change.

40. [ NH3 ] 2 Keq = [ N2 ] ● [ H2 ] 3 2 C6H6(g) + 15 O2(g) 12 CO2(g) + 6 H2O (g) [ CO2 ] 12● [ H2O ] 6 Keq = [ O2 ] 15● [ C6H6 ] 2 Examples 1) Write the equilibrium expressions for the following reactions.

41. [ NH3 ] 2 Keq = [ N2 ] ● [ H2 ] 3 [ 0.15 M ] 2 Keq = [ 0.5 M ] ● [ 0.26 M ] 3 2.6 Keq = Examples 2) At equilibrium the concentration of nitrogen is 0.50 M, hydrogen 0.26 M and ammonia 0.15 M. Calculate the equilibrium constant for the reaction.

42. [ NH3 ] 2 Keq = [ N2 ] ● [ H2 ] 3 [ 3.00 M ] 2 Keq = [1.00 M ] ● [ 4 M ] 3 0.141 Keq = Examples 3) In a 2.00 L reaction vessel there are 6.00 mol of ammonia, 2.00 mol of nitrogen and 8.00 mol of hydrogen at equilibrium. Calculate the equilibrium constant for the reaction.

43. [ NH3 ] 2 Keq = [ N2 ] ● [ H2 ] 3 n c = V [ x ] 2 11.1= n [1 M] ● [ 3M ] 3 17.3 = 1.0 L 17.3 M x= Examples 4) Ammonia is formed in a 1.00 L reaction vessel. If the equilibrium constant is 11.1 at a given temperature and the mixture contains one mole of nitrogen, calculate the number of moles of ammonia present at equilibrium. n = 17.3 mol

44. Examples • For the system A(g) + 2 B(g)→ 2 C(g) 0.500 moles of A and 1.00 mole of B were placed in 500mL reaction vessel. The equilibrium concentration of A was found to be 0.500 mol/L. Complete the following: • a. Use I.C.E to calculate the concentrations of B and C • A(g) + 2 B(g)→ 2 C(g) • I 1.0 M 2.0 M 0.0 M • C 0.5 M 1.0 M 1.0 M • E 0.5 M 1.0 M 1.0 M • b. What is the equilibrium constant expression and the value of k?

45. Examples Determine the [ ] of H+ a 0.500 mol/L CH3COOH (aq) solution. CH3COOH(aq)→ CH3COO -(aq) + H+(aq) I C E 0.500 mol/L 0 0 x x x (0.500 – x) x x Ka = [CH3COO-] [H+] [CH3COOH] X2 = (1.8 x 10-5)(0.500) X = 3.00 x 10-3 X = [H+] 1.8 x 10-5 = (x) (x) (.500 – x)

46. At equilibrium Keq > 1 Product is favored Keq = 1 Product and Reactant are equal Keq < 1 Reactant is favored P.419 #11 & 12

47. The Haber Process N2 (g) + 3H2 (g)2NH3 (g) As NH3 is formed, some of it back reacts and forms N2 and H2. This takes place until the amount consumed is equal to the amount produced.

48. iron catalyst 2. Compressor 3. Converter 200 atmospheres 450°C • Gases are mixed and scrubbed unreacted N2 and H2 are recycled N2 H2 NH3 4. Cooler N2 H2 NH3 to storage

50. 200 g (2 cup) 100 g (1 cup) Heterogeneous Systems • Substances are in different phases at equilibrium • i.e., solid and aqueous. Which solid is more concentrated? Concentration of a solid (and pure liquid) is always a constant